Lewis bonding theory

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The Lewis bonding theory is the first thing that comes to mind when discussing the visual representation of a chemical bond. The Lewis dot structure of a molecule shows how electrons from different atoms get involved in chemical bonding. A pictorial representation of any model is essential to grasp its understanding.

So, wait no more and dive into the article so we can discuss everything you need to know about the Lewis dot structures and the Lewis bonding theory.

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What is Lewis’s theory of chemical bonding – Definition

Lewis dot structure is the most simplified representation of the valence electrons present in a chemically bonded molecule. This concept was introduced by an American chemist Gilbert Newton Lewis in 1916. Irving Langmuir further endorsed the idea in 1920. As per Lewis’s theory, only valence (outer shell) electrons are involved in chemical bonding. A new chemical bond is formed between two atoms by the complete transference or mutual sharing of valence electrons such that both achieve a stable noble gas electronic configuration.

Importance of the Lewis model in covalent bond formation

  • Lewis bonding theory defines the formal charges present on the bonded atoms in a molecule and, consequently, the overall charge present on it. Formal charges help predict the stability of a molecule. The fewer the formal charges on the bonded atoms in a Lewis structure, the more stable it is.
  • It is important to understand concepts such as electron affinity and electronegativity before drawing the correct Lewis structure of a molecule. Electronegativity determines the polarity of a molecule. In this way, the Lewis model is indirectly related to the polarity of a covalent bond.
  • We can also determine the resonance present in a molecule by showing the movement of electrons from one position to another in its Lewis structure.

What is a Lewis dot structure?

The Lewis dot structure, also known as an electron dot diagram, shows how the valence electrons are arranged in individual atoms. In addition, it also shows how valence electrons of two different atoms get involved in forming a new chemical bond.

In a Lewis dot structure:

  • Atoms are shown by their atomic symbols.
  • The dots represent electrons.
  • A single straight line represents a chemical bond or a bond pair containing 2 electrons.
  • Two straight lines show a double bond, while three straight lines denote a triple covalent bond.
  • An unbonded or lone pair of electrons present in an atom are shown as two dots placed together.
  • A total of 8 dots around an atomic symbol denotes a complete octet.

Lewis dot representation for an ionic compound

An ionic bond is formed by the opposite charge attraction between a highly electronegative non-metal atom and an electropositive metal atom.

 For example, let’s see how to draw the Lewis dot structure of MgF2, an ionic compound.

The electronic configuration of Mg is 1s2 2s2 2p6 3s2.

The Mg atom loses two electrons from its outermost shell (n= 3) to form Mg2+.

As the Lewis structure only shows valence electrons, thus there are no dots to show in the Lewis structure of Mg2+. Hence it is simply shown as:

In contrast, an F-atom gains 1 electron in its valence shell to form F.

The electronic configuration of F is 1s2 2s2 2p5. The gain of an electron denotes a total of 8 dots around the F symbol in the Lewis structure.

F Lewis structure is multiplied by 2 to represent that two F-atoms gain the electrons lost by one Mg-atom to form MgF2.

Other examples for drawing the Lewis structures of ionic compounds are shown below.

You may learn more about ionic compounds through the electronic theory of chemical bonding.

Lewis dot representation for a covalently bonded molecule

The Lewis dot structures are most important for representing covalently bonded molecules, i.e., the molecules formed by the mutual sharing of valence electrons between corresponding atoms.

Let’s see as an example how to draw the Lewis structure for O2 using the step-by-step guide below.

Step I: Count the total valence electrons present in the molecule.

O2 is made up of two identical O-atoms.

The electronic configuration of an O-atom is 1s2 2s2 2p4.

There are a total of 6 valence electrons in an O-atom. 2 O-atoms denote 2 x 6 = 12 valence electrons.

Step II: Represent the covalent bond using single straight lines.

An O2 molecule consists of a double covalent bond, i.e., both the O-atoms are joined to each other using 2 straight lines.

Each straight line represents a bond pair (2 electrons).

Step III: Represent the remaining electrons as lone pairs.

4 valence electrons consumed out of the 12 initially available leaves behind 12-4 = 8 valence electrons, i.e., 2 lone pairs on each O-atom.

So, the final Lewis structure for O2 is shown below. Each O-atom has a complete octet in this molecule.

The formal charge on an atom = Total valence electrons in a free atom – Unbonded electrons – ½ bonding electrons

Formal charges on both the O-atoms in O2 = 6 – 4 –4/2 = 6– 4 –2 = 6– 6= 0.

Zero formal charges are present on both atoms; thus, there is no overall charge present on the O2 Lewis structure, and it is a stable molecule.

The O2 molecule comprises two identical O-atoms with the same electronegativity value (E.N = 3.44). So the O=O electron cloud stays equally shared between the two O-atoms. Consequently, O2 is a non-polar molecule.

Similarly, you can draw the Lewis structures for more complicated molecules such as H2O, SO2, NH3, etc., as shown below.

All the molecules drawn above are polar.

As their Lewis structures show, these molecules are made up of dissimilar atoms. A specific electronegativity difference exists between the bonded atoms in each of the O-H, S=O and N-H bonds, denoting bond polarity. The highly electronegative atom more strongly attracts the shared electron cloud from each bond pair which leads to a non-uniform charge distribution. The unsymmetrical molecular shape further enhances the polarity effect and makes the molecule polar overall.  

The molecular shape concept is out of the scope of the Lewis theory. So you can find a detailed discussion on it in the VSEPR theory of chemical bonding.

Using Lewis structures, we can show the resonance present in a molecule, for example, the SO3 molecule consists of the following three resonance forms.

The pi-bonded and un-bonded electrons are delocalized, so they keep moving from one position to another on the molecule. The actual SO3 Lewis structure is thus an average or hybrid of all the corresponding resonance structures.

We have also provided you with the Lewis structures of some basic organic compounds which are essential for you to learn if you are a chemistry student.

Limitations of Lewis theory and future models

The Lewis bonding model is a simple approach, primarily helpful for representing ionic and covalent bonds. It is merely a qualitative approach to explain the bonding phenomenon.

More extensive models, such as the VSEPR theory and VBT, followed the Lewis theory to discuss other important features of chemical bonding, such as orbital overlap, molecular shape, size, geometry, bond energies and bond angles etc., about which there was no information given by the Lewis bonding theory.

You may also like the following:


1. Logan, S. R. 2001. ‘The Role of Lewis Structures in Teaching Covalent Bonding’, Journal of Chemical Education, 78: 1457.

2. Shah, Nilpa. 2019. “Lewis theory of bonding ” In. LibreTexts Chemistry.  

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